Question

Why does the electrolysis of an aqueous sodium chloride solution produce hydrogen gas at the cathode?

Short answer

Hydrogen gas is produced at the cathode during the electrolysis of aqueous sodium chloride solution because water is reduced there instead of sodium due to its lower reduction potential.

Step-by-step solution

- Understanding the process of electrolysis

Electrolysis is the process that uses an electric current to drive a non-spontaneous chemical reaction. When aqueous sodium chloride (NaCl) is electrolyzed, the resulting chemical reactions at the electrodes produce chlorine gas at the anode and hydrogen gas at the cathode.

- Identifying the reactions at the electrodes

At the cathode, water is reduced instead of sodium because it has a lower reduction potential. The reduction reaction for water is: \( 2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq) \). This reaction produces hydrogen gas.

- Understanding why sodium is not reduced

Sodium has a more negative reduction potential than water, which means that it is less likely to gain electrons than water. Thus, under standard conditions, water is reduced to hydrogen gas rather than sodium metal being deposited at the cathode.

- Reiterating the production of hydrogen gas at the cathode

We confirm that during the electrolysis of aqueous NaCl, hydrogen ions from water molecules at the cathode are reduced to hydrogen gas, and this is observed as bubbles rising from the cathode.

Key concepts

Understanding Chemical Reactions in Electrolysis

Electrolysis is a fascinating application of chemistry that allows us to induce a chemical change via an electric current. In simple terms, think of it as using electricity to break apart compounds or create new substances. During the electrolysis of an aqueous sodium chloride solution, also known as table salt dissolved in water, we see this principle in action.

Chemical bonds are influenced by the presence of electric current, causing elements or compounds to either lose or gain electrons - a process we call oxidation or reduction, respectively. In the case of sodium chloride in water, chlorine is oxidized at the anode, producing chlorine gas, while at the cathode, water undergoes reduction to form hydrogen gas and hydroxide ions. In essence, we have two key chemical reactions happening side by side, each with its own importance in the overall process of electrolysis.

The Role of Reduction Potentials in Electrolysis

For students diving into the world of electrolysis, reduction potential is a critical concept to grasp. It's like a 'tug-of-war' between atoms and ions, where electrons are the 'rope' and the winner is the one that has an easier time pulling these electrons towards itself. This 'ease' is quantified by the reduction potential.

To understand why water is reduced to hydrogen at the cathode instead of sodium, we consider the reduction potentials of both reactions. The negative value of sodium's reduction potential means it's less inclined to gain electrons compared to water, which has a more positive reduction potential. Because of this, when you're running an electric current through our saltwater solution, water 'wins', and hydrogen gas is produced. It's like betting on the stronger team; water is simply more likely to take those electrons and transform into hydrogen gas.

Production of Hydrogen Gas at the Cathode

The production of hydrogen gas is not just a theoretical outcome; it's a visible and tangible result of the electrolysis of sodium chloride solution. If you were to observe the process, you'd see bubbles emerging at the cathode, a clear sign that a gas is being produced. This is hydrogen, the lightest and most abundant element in the universe, and it forms when the two hydrogen ions (from water molecules) receive electrons during the reduction process.

Why bother producing hydrogen gas, you might ask? Hydrogen has a vast array of uses, from serving as a clean energy source to being a fundamental element in various industrial processes. Therefore, the action you're seeing at the cathode is not just a neat chemical trick—it's a process with significant real-world applications.